Fuel 111 (2013) 40–47

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Conversion of CO2 into mineral carbonates using a regenerable bufferto control solution pH

0016-2361/$ - see front matter � 2013 Elsevier Ltd. All rights reserved.http://dx.doi.org/10.1016/j.fuel.2013.04.033

⇑ Corresponding author. Tel.: +61 733653977; fax: +61 733654199.E-mail address: karen.steel@uq.edu.au (K.M. Steel).

Karen M. Steel ⇑, Kimia Alizadehhesari, Reydick D. Balucan, Bruno BašicSchool of Chemical Engineering, University of Queensland, St. Lucia, Queensland 4072, Australia

h i g h l i g h t s

� Use of a regenerable buffer (tertiary amine) is studied to enable mineral carbonation.� Tertiary amines complex protons to give a pH of 8.2 which enables MgCO3 precipitation.� Increasing temperature from 18 to 85 �C decreases pH by 2.5 pH units.� Higher temperatures >85 �C might enable low pHs needed for Mg–silicate dissolution.

a r t i c l e i n f o

Article history:Received 25 October 2012Received in revised form 14 March 2013Accepted 16 April 2013Available online 30 April 2013

Keywords:Mineral carbonationCO2 sequestrationCO2 mineralisationRegenerable buffer

a b s t r a c t

The barrier that is currently stalling the rapid conversion of magnesium silicate deposits into magnesiumcarbonate as method for storing CO2 is considered to be the difference in pH needed for magnesium dis-solution from the silicate and magnesium precipitation as the carbonate, whereby rapid dissolutionrequires a low pH of around 1 while rapid precipitation requires a considerably higher pH of around 8.This paper investigates a novel concept which is to use a tertiary amine to bind with protons and raisethe pH to around 8 and to then regenerate the amine through the use of heat due to the strength ofthe amine–proton bond decreasing with increasing temperature. This approach provides the low pHand high temperature that is needed for Mg dissolution and the high pH need for carbonate precipitation.The amine can be thought of as a regenerable buffer.

Dissolution of Mg from serpentine has been found to be favourable with a solids to solution volume ofmore than 50 g/L to enable a low pH, and with temperatures close to the boiling point of the solution. ThepH needed for magnesium carbonate precipitation was found to be approximately 8.2. Both triethyl-amine and tripropylamine were found to be capable of achieving this at 18 �C. Yields of around 20–40 wt.% carbonate were achieved using residence times of approximately 1 h. The pH swing for the ter-tiary amines was found to be approximately 2.5 pH units between 5 and 85 �C, suggesting that an aminecapable of achieving a pH of 8.2 at low temperature generates a pH of 5.7 in solution when heated to85 �C. Further work will examine whether the lower pH values needed for serpentine dissolution canbe achieved by heating the protonated amine to higher temperatures.

� 2013 Elsevier Ltd. All rights reserved.

1. Introduction

This paper investigates a novel concept to enable the conversionof CO2 into mineral carbonates using an aqueous route. The barrierthat is currently stalling conversion is pH control. The first part ofthis introduction outlines why pH is critical, the second part out-lines what work has been done in the field of converting CO2 intomineral carbonates using an aqueous route, and the third partintroduces the novel concept behind this study.

1.1. Thermodynamic modelling of the carbonate system

In order for a mineral carbonate, such as MgCO3, to form theconcentrations of CO2�

3 and Mg2+ in solution must be high enoughfor Eq. (1) to be satisfied [1].

½Mg2þ�½CO2�3 � > 3:46� 10�8 ð1Þ

The solubility of CO2 in aqueous solution and the dissociation ofcarbonic acid to form bicarbonate and carbonate anions can be de-scribed by Eqs. (2)–(5) [1–4].

PPCO2 ðatmÞ ¼ 29:36½H2CO3� ð2Þ

½HCO�3 � ¼ ½H2CO3�=ð2:249� 106½Hþ�Þ ð3Þ

CO2 partial pressure (atm)10-1110-10 10-9 10-8 10-7 10-6 10-5 10-4 10-3 10-2 10-1 100 101 102

Con

cent

ratio

n (M

)

10-1210-1110-1010-910-810-710-610-510-410-310-210-1100101

H2CO3

HCO3-

CO32-

Fig. 2. Equilibrium concentration of H2CO3, HCO�3 and CO2�3 in water as a function

of CO2 partial pressure (derived from Eqs. (2)–(5)).

pH1 2 3 4 5 6 7 8 9 10 11

CO

32- c

once

ntra

tion

(M)

10-1510-1410-1310-1210-1110-1010-910-810-710-610-510-410-310-210-1100101

0.1atm 1 atm 10 atm 100 atm

Fig. 3. Equilibrium concentration of CO2�3 as a function of pH for CO2 partial

pressures of 0.1, 1, 10 and 100 atm.

K.M. Steel et al. / Fuel 111 (2013) 40–47 41

½CO2�3 � ¼ ½HCO�3 �=ð2:133� 1010½Hþ�Þ ð4Þ

½Hþ� ¼ ½HCO�3 � þ 2½CO2�3 � ð5Þ

where 29.36 is the Henry’s law constant (dm3 atm mol�1) for CO2 inwater at 25 �C [2]. Here, [H2CO3] includes both dissolved molecularCO2 and molecular H2CO3. Eqs. (3) and (4) shows that CO2�

3 ionsonly start to become significant in solution when the pH is aboveabout 8 (see Fig. 1). It could be thought that if the partial pressureof CO2 was raised the concentration of all species in solution(H2CO3, HCO�3 and CO2�

3 ) would increase, and therefore a raised con-centration of CO2�

3 would result. However, solving Eqs. (2)–(5) forvarious partial pressures of CO2 (PPCO2 ) shows that the concentra-tion of CO2�

3 in solution stays very low at around 5 � 10�11 M de-spite the partial pressure of CO2 rising to 50 atm (see Fig. 2). Thisis due to the H+ ions, forming from the dissociation of H2CO3, alwayssuppressing the formation of CO2�

3 (see Eqs. (3) and (4)). Only theconcentration of HCO�3 becomes appreciable at high CO2 partialpressures.

Given the solubility limit of MgCl2 in solution (55 g/100 ml), themaximum possible concentration of Mg2+ is approximately6 M Mg2+ and the corresponding carbonate concentration requiredto precipitate MgCO3 is 5 � 10�9 M. Given that 6 M Mg2+ is difficultto achieve as it is close to the solubility limit, a more realistic con-centration of 0.6 M Mg2+ would require a CO2�

3 concentration of5 � 10�8 M CO2�

3 . Assuming a partial pressure of CO2 of 1 atm,Eqs. (2)–(4) can be solved to establish a relationship between car-bonate concentration and pH. Fig. 3 shows this relationship andshows that when the carbonate concentration is 5 � 10�8 M theequilibrium pH is approximately 5.5. This means that in order toprecipitate MgCO3 the pH needs to exceed 5.5.

1.2. Work to date on converting CO2 into mineral carbonates using theaqueous route

Bond and co-workers [5–7] looked at the potential of using anenzyme, called carbonic anhydrase, to convert power station CO2

into CaCO3. Coral reefs use carbonic anhydrase to assist in the pro-duction of CO2�

3 . The enzyme was found to catalyse the formationof CO2�

3 , however, although the catalyst enables the rapid forma-tion of CO2�

3 , there is also a simultaneous rapid drop in solutionpH to approximately 4, which prevents the precipitation of carbon-ates. In the work of Bond et al. [5–7] and in recent work by Mirja-fari et al. [8], Ozdemir [9] and Rayalu and co-workers [10,11],precipitation of carbonates was only possible if a buffer was addedto the solution in order to complex H+ ions and keep the pH high.

pH1 2 3 4 5 6 7 8 9 10 11 12 13 14

mol

frac

tion

0.0

0.2

0.4

0.6

0.8

1.0

H2 CO3

HCO3-

CO32-

Fig. 1. Equilibrium distribution of H2CO3, HCO�3 and CO2�3 species in solution.

The buffer used was tris(hydroxymethyl)aminomethane or ‘Tris’,which is commonly used to maintain pH at around 8. This buffercould not be continuously used in a large scale process, as it cannotbe regenerated, which presents a major obstacle to the furtherdevelopment of this approach to sequestering CO2. Coral reefs havethe ocean as a giant buffer for the H+ that they generate. It isworthwhile to note that Mirjafari et al. [8] found that calcium car-bonate did not precipitate when they used carbonic anhydrasewith no buffer, but found precipitation to take place when theyused the buffer with no carbonic anhydrase.

Instead of converting salt solutions into carbonates, researchershave also looked at converting Mg–silicate minerals into carbon-ates in order to use the neutralising capacity of the mineral [12].O’Connor and co-workers [13,14] found they could convert 34%of serpentine (Mg3Si2O5(OH)4) to MgCO3 using aqueous CO2 at115 atm, 185 �C and a residence time of 24 h. The pH generatedin solution would be around 3. The high temperature would beassisting the kinetics of Mg dissolution. Presumably the pH in-creases as dissolution proceeds and this aids precipitation of car-bonate. However, as the pH increases the dissolution rate of Mgwould drop to negligible levels. The silica left behind and carbon-ate forming would also hinder further Mg dissolution/carbonation.O’Connor et al. found that artificially adding NaHCO3 to shift theequilibria in favour of a higher CO2�

3 to H+ ratio aided carbonation.

42 K.M. Steel et al. / Fuel 111 (2013) 40–47

It follows that the best way to enable carbonation might be tohave a two stage process where the first stage is optimised forMg dissolution (high temperature and low pH) and the second isoptimised for Mg carbonation (high pH). This need has been recog-nised as pH swing [15]. pH swing requires the removal of H+ ionsfrom solution, which is usually achieved with soluble oxides/hydroxides, however, these come from the calcination of carbon-ates which is obviously not possible.

1.3. Novel concept for converting CO2 into mineral carbonates usingthe aqueous route

This paper investigates the use of weakly basic tertiary aminesfor complexing H+ ions generated when CO2 is added to solutionand therefore enabling carbonate precipitation from a variety ofsalt solutions. The acid dissociation constant for protonated ter-tiary amines varies with temperature and so the approach is touse ‘‘temperature swing’’ to regenerate the amine, whereby theloaded amine is heated to strip the acid off. This concept has re-cently been patented [16].

Conventional CO2 capture technologies involve absorbing CO2

into a mixture of primary and tertiary amines, including monoeth-anolamine (MEA) and methyldiethanolamine (MDEA) respectively.The reason for this is that while the primary amine forms a strongcarbamate bond and therefore enables a high CO2 loading due tothe strength of the bond considerable energy is needed to breakit and regenerate the MEA (approximately 3–5 MJ/kg CO2 [17]).In order to reduce the energy load tertiary amines are blended. Ter-tiary amines do not bond with CO2 but rather strip the solution ofH+ ions thus driving the formation of HCO�3 and CO2�

3 ions in solu-tion. The loading of CO2 in solution as these ions is much lower butthe energy needed to regenerate the MDEA is much less, such thatthe total energy needed to regenerate the MEA/MDEA mixture isaround 1–3 MJ/kg CO2 [17]. The introduction of MDEA means thattaller towers are needed for a given separation efficiency and so abalance between capital and operating (energy) cost must bestruck.

The idea put forward here is to use tertiary amines alone to con-vert CO2 into mineral carbonates. It is thought that the loading ofCO2 can be higher than that for conventional CO2 capture usingMDEA because the CO2�

3 ions that form leave the solution as solidcarbonate, thus providing a stronger driving force for the capture ofCO2 into solution.

The advantage of this concept for CO2 sequestration is twofold.Firstly, the energy needed for the process might be considerablylower than that needed for conventional CO2 capture because lessenergy is needed to regenerate the amine and compression of theCO2 (needed for storage) is not necessary. Secondly, the CO2 wouldbe locked up in a mineral form that is known to be stable for mil-lions of years, which is an important consideration given the scaleof CO2 that needs to be stored.

Fig. 4. Simplified block flow diagram of the

Fig. 4 shows a simplified diagram of the novel concept that isbeing explored. In stage 1, acid loaded amine is heated to�100 �C and contacted with a Mg silicate such as serpentine((Mg, Fe)3Si2O5(OH)4). At the high temperature the acid (HCl) dis-sociates from the amine, thereby providing a low pH capable ofdissolving Mg out of the serpentine. The Mg depleted serpentineis separated by density and/or filtration and the solution, contain-ing MgCl2 and regenerated amine, passes to stage 2. In stage 2, thesolution is cooled and the flue gas containing CO2 is spargedthrough. At the low temperature acid generated by the dissociationof H2CO3 plus excess acid generated in stage 1 are complexed bythe tertiary amine which causes the pH to rise and consequentlythe CO2�

3 concentration to rise to a level that is sufficient to begininteracting with Mg2+ and precipitating MgCO3. The MgCO3 is sep-arated by density and/or filtration and the solution, containing acidloaded amine is recycled to stage 1.

The overall reaction taking place in stage 1 is as follows:

Mg3Si2O5ðOHÞ4 þ 6R3NHCl! 3MgCl2 þ 6R3Nþ 5H2O

þ 2SiO2 ðR1Þ

And the overall reaction taking place in stage 2 is as follows:

MgCl2 þ CO2 þH2Oþ 2R3N ¡ MgCO3 þ 2R3NHCl ðR2Þ

The concept shown in Fig. 4 can be used in a variety of differentways by a variety of different industries, and is not locked into oneindustrial application.

The concept could be used for the treatment of magnesium sil-icate deposits as described above. It is worthwhile to note thatmany magnesium silicate deposits contain significant levels ofvaluable heavy metals such as Ni, called Ni laterites, and the pro-cess could therefore have the dual operation of Ni extraction com-bined with CO2 sequestration. It is known that treating Ni lateriteswith acid to dissolve the Mg enables a greater amount of Ni to beextracted into solution.

Secondly, the concept could be used for salt mining operations.Chloride and sulphate salt deposits are mined for KCl or K2SO4 tobe used as fertiliser. By using solution mining the salt comes uphot and is cooled to separate the potassium salts. The refuse saltrepresents a waste that is generally deposited back into the re-serve. The salt solution could be processed using the sequestrationtechnology shown in Fig. 4 to form a mixture of carbonates andbicarbonates that are deposited back into the reserve for longtermCO2 storage. The heat that must be taken out of the solution as itcomes to the surface could be used to provide the heat neededfor stage 1. Using the concept for salt solutions means that a by-product of the process is acid, either hydrochloric or sulphuric acid.Therefore, the scale of the operation would need to be matchedwith HCl or H2SO4 needs in the oil, chemical and mineral sectors.

This paper presents our work to date on this novel concept.

proposed CO2 sequestration technology.

K.M. Steel et al. / Fuel 111 (2013) 40–47 43

2. Experimental

2.1. Serpentine sample

The serpentinite sample (Mg3Si2O5(OH)4) used for this studywas obtained from a naturally occurring deposit in northernQueensland, Australia. It was initially ground by hand in a pestleand mortar and then in a laboratory attrition mill. The ground ser-pentinite was sieved with ASTM standard sieves to obtain particleswith a diameter of <57 lm. Australian Laboratory Services (ALSs)performed the elemental analysis via alkali fusion, acid digestionand inductively coupled plasma-atomic emission spectroscopy(ICP-AES) of the resulting solution. The loss on ignition at1000 �C (LOI1000) was also performed using a TGA furnace. The re-sults of ALS’s analysis based on their method ME-ICP85 (Silicatesby Fusion, ICP-AES) and ME-GRA05 (H2O/LOI by TGA Furnace)are summarised in Table 1.

Mineral composition was first probed via X-ray diffraction anal-ysis using a PANanalytical XPERT-PRO diffractometer with Cu Katarget (k = 0.15406 nm) at room temperature. Measurements weremade in a step scan mode (0.1�/step) over the 2h range of 10–90�.Phase matching of the X-ray powder diffraction (XRPD) pattern ofthe serpentinite sample against the International Centre for Dif-fraction Data (ICDD) database suggested antigorite to be the pri-mary serpentine phase present. A calibrated TA InstrumentSDTQ600 Thermogravimetric analyser-differential scanning calo-rimeter provided further mineral characterisation via thermogravi-metry–derivative thermogravimetric analysis (TGA–DTG).Replicate runs were obtained for 10 mg of the �53 lm samplesusing alumina crucibles and heated from 30 �C to 1000 �C at a heat-ing rate (b) of 10 �C min�1. A 10-min isothermal stage was em-ployed at 110 �C to determine the moisture content of thematerial and was found to contain 1.0 wt.%. The total mass lossof the dry sample (Dm105–850 �C) was determined as11.5 ± 0.01 wt.%, which is in fair agreement with the analysis madeby ALS (11.7 wt.%).

Fig. 5 shows the TGA–DTG profile of the sample, where thecharacteristic serpentine doublet comprising the DTG temperatureshoulder, Tsh, showing at 597 �C and the peak temperature, Tp, at718 �C. Thermal analysis suggests that this particular serpentinitesample is fully serpentinized and contains antigorite as well aslizardite (antigorite + lizardite). The antigorite component displaysits DTG peak temperature, Tp1ATG at 718 �C and its diagnostic peak,Tp2ATG at 747 �C. The shoulder at 701 �C is thought to indicate thelizardite component, Tp1LIZ. Based on XRPD and TGA-DTG analysis,we then refer to this sample as ‘‘serpentinite’’, rather than antigor-ite as this specimen also contains lizardite.

2.2. Magnesium dissolution

Magnesium dissolution experiments were carried out using ARgrade HCl (37 wt.%) and Millipore water. 0.5 g of sample and acidsolution was mixed in a 250 ml spherical flat-bottom flaskmounted on a magnetic stirrer/hotplate and equipped with a coldwater condenser. The effects of temperature, concentration of HCl,

Table 1Chemical composition of the serpentinite sample.

wt.%, ±0.01MgO* SiO2

* Fe2O3* Al2O3

* CaO* Ni* MnO* K2O* LOI1000**

39.3 44.1 7.38 1.03 0.35 0.24 0.10 0.04 11.7

* Values obtained via inductively coupled plasma-atomic emission spectroscopy onfused samples after acid digestion. Based on ALS’s ME-ICP85.** Value obtained by heating the moisture free sample to 1000 �C using TGA fur-nace. Based on ALS’s ME-GRA05.

residence time and acid solution volume were investigated. At theduration of the experiment, the residue was vacuum filtered, driedovernight and weighed. The pH of the filtrate solutions was mea-sured using a pH electrode. The solid remaining was analysed byALS using the procedure described above to determine the extentof Mg dissolution.

2.3. Carbonate precipitation

Carbonate precipitation experiments were carried out using a250 ml Erlenmeyer flask open to the atmosphere. Food grade CO2

was injected into the solution through a sparger containing fiveholes of 2 mm diameter each. Pressure was regulated at 1.4 barand flow was set at approximately 1 L/min using a rotameter.While sparging the solution with CO2, tertiary amine was addeddropwise via a burette while simultaneously measuring pH via apH electrode. The amines investigated were simple straight chaintrialkylamines with increasing chain length, i.e. triethylamine, tri-propylamine, tributylamine, tripentylamine, trihexylamine. Thesolution was observed for the onset of precipitation. If a precipitateform, it was filtered through Whatman No. 1 filter paper, dried andweighed. Precipitated solids were analysed by XRD and ICP-AES forcompound and elemental determinations, as described above.

2.4. Amine regeneration

The ability of the amines to be regenerated and liberate boundacid was investigated via a series of titrations at various tempera-tures whereby the amines were added to a standardised solution of0.1 M HCl (10 ml). The HCl was placed in a 3-neck round bottomflask which was immersed in a water bath to control temperature.A condenser was fitted vertically to the middle neck and a ther-mometer and pH electrode were inserted and sealed through eachof the side necks. Amine was added through the opening at the topof the condenser using an automatic pipette. For experiments per-formed at 5 �C ice was added to the water bath. The tertiary aminesstudied were triethylamine, tripropylamine, tributylamine andtripentylamine.

3. Results and discussion

3.1. Magnesium dissolution

The effect of HCl concentration on the dissolution of Mg fromserpentine is shown in Fig. 6. All percentages are weight percent-ages. The residence time used was 3 h and the temperature wasthe boiling temperature of the solution (�100 �C). The stoichiome-tric amount of acid needed to dissolve all of the Mg according to R1is approximately 0.12 M which gives approximately 40% extrac-tion, while a plateau of approximately 65% extraction is reachedat around 0.5 M HCl. Fe and Al were also found to dissolve withsimilar extraction levels to those of Mg, showing that the elementsdo not appear to dissolve selectively with respect to HClconcentration.

Because it is desirable to not have excess acid in solution tominimise the energy needed for amine regeneration a compromisebetween Mg extraction and solution pH must be struck. Fig. 6shows the pH change as a function of HCl concentration. Withtwice the stoichiometric amount needed, the pH of the spent solu-tion is still �1.

The effect of residence time on dissolution of Mg in 0.25 M HClis shown in Fig. 7. Equilibrium is reached after approximately 3 h.The effect of temperature on the dissolution of Mg in 0.25 M HCl isshown in Fig. 8. For temperatures less than 50 �C only a smallamount of Mg dissolves. At temperatures higher than 50 �C a linear

Fig. 5. The TGA–DTG profile of the serpentinite sample used in this study with the characteristic peaks indicated.

Fig. 6. Effect of HCl concentration on the dissolution of Mg and final pH (residencetime 3 h, �100 �C, 0.5 g, 100 ml).

Fig. 7. Effect of residence time on the dissolution of Mg (0.25 M HCl, solution and�100 �C, 0.5 g, 100 ml).

Fig. 8. Effect of temperature on the dissolution of Mg (0.25 M HCl, residence time3 h, 0.5 g, 100 ml).

Fig. 9. Effect of solution volume on the dissolution of Mg (0.025 mols HCl, residencetime 3 h, 0.5 g, �100 �C).

44 K.M. Steel et al. / Fuel 111 (2013) 40–47

increase in Mg dissolution with respect to temperature is obtained,reaching approximately 65% at �100 �C and suggesting that attemperatures above 100 �C higher extraction efficiencies mightbe achieved. Extrapolation suggests that close to 100% extractionmight be achieved at 140 �C. Experiments using a pressurised ves-

sel to enable higher temperatures above 100 �C are recommendedto confirm the extrapolated trend shown.

As presented in the introduction, it is desirable to have a highMg2+ concentration in solution as this reduces the concentrationof CO2�

3 needed for MgCO3 precipitation. Experiments were per-

Fig. 10. Titration curve for triethylamine against 0.1 M HCl at 18 �C and 90 �Cgenerated from pKa data obtained from literature and at 18 �C generated fromexperiment.

K.M. Steel et al. / Fuel 111 (2013) 40–47 45

formed keeping the amount of HCl the same and decreasing thesolution volume from 100 ml. Twice the stoichiometric amountneeded (0.024 mols) was chosen for the amount of HCl. Fig. 9shows the effect of reducing the solution volume down to 10 ml.The extraction increases as the solution volume decreases reachingapproximately 85% with only 10 ml of solution. The concentrationof Mg in solution is approximately 0.35 M.

This work has shown the importance of both acid concentrationand temperature on the dissolution of Mg. It is recommended tooperate with a solids to solution volume of more than 50 g/L forthe extraction stage, a temperature close to the boiling tempera-ture of the solution or higher if using a pressurised vessel, a resi-dence time of 3 h and concentration no more than twice thestoichiometric amount needed for reaction. These conditions en-able high extractions of Mg approaching 100%.

3.2. Carbonate precipitation

The extract solution from the experiment with 10 ml solutionvolume shown in Fig. 9 was used for carbonation. Tripropylaminewas added dropwise. As the pH increased to approximately 5 alight brown precipitate formed which was found from elementalanalysis to contain around 20.1 wt.% Fe, 10.8 wt.% Si and 5.5 wt.%Al and only 0.2 wt.% Mg. This product comes from the hydrolysisof Fe, Si and Al which dissolved during serpentine dissolution.The leaching studies had shown that the elements dissolved simul-taneously with Mg. After removing this precipitate further addi-tions of TPA raised the pH to approximately 8 at which point awhite precipitate formed. This precipitate began forming withina few minutes. After bubbling CO2 for 45 min, the precipitate wasrecovered by filtration, dried and analysed. The weight was0.20 g. XRD analysis indicated the formation of nesquehonite(MgCO3�3H2O). Elemental analysis showed that the purity washigh with a composition of 18.8 wt.% Mg, 0.2 wt.% Ca, 0.06 wt.%Fe and 0.01 wt.% each of Al and Si. The yield was 29.1 wt.% (i.e.29.1 wt.% of the Mg extracted into solution was converted to theMgCO3 precipitate). The remaining Mg would be recycled to thefirst stage of the process.

In order to study MgCO3 more precisely and the pH level neededfor carbonate precipitation without the hindrance of other dis-solved elements model compound work was performed usingMg(OH)2. 0.126 g of Mg(OH)2 was dissolved in 50 ml of 0.0965 MHCl such that the acid was slightly in excess and gave a final pHof 2.08 after the Mg(OH)2 had completely dissolved. To this solu-tion 5 � 10�4 mols of TPA was added, which is the amount neededto neutralise the excess acid. The pH rose to 9.59. CO2 was thenbubbled through the solution and after 5 min the pH had decreasedand stabilised at 4.62 and no precipitate had formed. A furtheraddition of TPA was made (0.006 mols) and the pH stabilised atapproximately 8.27. This addition of TPA is 1.5 times that neededfor reaction 2. Over the next half an hour, CO2 was continually bub-bled through the solution. It was found that TPA need to be contin-ually added dropwise in order to maintain the pH at a level above8. The total amount of TPA added (neglecting the initial 5 � 10�4 -mols) was 0.021 mols and the final pH was 8.43. The solution wasfiltered and the solid recovered and air dried. The weight of the so-lid was 0.051 g and as found above, analysis indicated nesqueho-nite (MgCO3�3H2O) to be the primary phase present. The yieldwas approximately 17% and the amine used was 5 times in excess.

The above experiment was repeated with slower additions ofTPA. A total amount of 0.0085 mols (twice excess) was added overa period of 1 h. The mass of solid recovered was 0.073 g (dried)which is a yield of 24 wt.%.

The above experiment was repeated with triethylamine (TEA).0.164 g of Mg(OH)2 was dissolved in 50 ml of 0.112 M HCl, whichis the stoichiometric amount needed for complete dissolution.

CO2 was bubbled which decreased the pH to 5.38. 0.0072 mols ofTEA was added and the pH increased to 10.16 and simultaneouslythe solution became milky with precipitation. As CO2 addition con-tinued the pH decreased to 7.27 even though the amount of TEAadded was 30% above that needed for reaction 2. A further additionof 0.0036 mols of TEA increased the pH to 9.32 initially but then itstabilised at 8.23. The solution was filtered to recover the solid,which had a mass of 0.133 g and therefore a yield of 34 wt.%.

Tests with both tributylamine and tripentylamine did not yieldprecipitates, which is thought to be due to the pH generated by theamines not being high enough. It is possible that decreasing thetemperature would enable precipitation to take place with theseamines.

These experiments have shown that the pH needed for magne-sium carbonate precipitation is approximately 8.2 and that trieth-ylamine and tripropylamine are capable of achieving this. Anexcess of amine has been found to be necessary to maintain thepH of 8 while CO2 is bubbled through the solution. So far yieldsof around 20–40 wt.% have been achieved. The reason why precip-itation did not occur at lower pH levels, such as the pH level of 5.5predicted from the theoretical modelling work reported in theintroduction, is thought to be due to the kinetics being too slow be-low 8.2. It was found with TEA that precipitation was within sec-onds when the pH was 10.

Further experiments will investigate the kinetics of carbonateprecipitation by sampling periodically, particularly during theearly stages.

3.3. Amine regeneration

Based on the serpentine dissolution work, to achieve dissolu-tion of Mg under reasonable conditions of a residence time of lessthan an hour and reaction time of 100–150 �C, the pH of the solu-tion needs to be approximately less than 1. Based on the carbon-ation work, the pH of the Mg rich solution needs to be raised toapproximately 8.2. To examine the ability of tertiary amines toreversibly enable this change a series of titrations were performed.The dissociation constant for triethylamine at various tempera-tures has been published by Hamborg and Versteeg [18], wherebythe pKa is 10.89 at 18 �C and decreases to about 9.17 at 90 �C.Fig. 10 shows these constants converted into a titration curvewhereby the amine is being added to 0.1 M HCl. Our own titrationpoints are also shown for comparison. The titration curve is ex-pressed this way as it mimics the real system. Total amine essen-

Fig. 11. Titration curve for tripropylamine against 0.1 M HCl at 18 and 85 �C.

Fig. 12. Titration curve for tributylamine against 0.1 M HCl at 5, 18 and 85 �C.

Fig. 13. Titration curve for tripentylamine against 0.1 M HCl at 5, 18 and 85 �C.

Table 2Calculated pKa values for tertiary amines at 5, 18 and 85 �C.

Amine pKa

5 �C 18 �C 85 �C

Tripropylamine nd 9.47 7.32Tributylamine 9.84 8.44 6.32Tripentylamine 6.81 5.80 3.98

nd: Not determined.

Fig. 14. Expected trend for pKa as a function of temperature for tributylamine andtripentylamine with lines of best fit, and pKa line for triethylamine as derived fromHamborg and Versteeg [18].

46 K.M. Steel et al. / Fuel 111 (2013) 40–47

tially means the amount of amine added expressed as a concentra-tion. At 20 �C the pH rises to approximately 11 while at 90 �C itrises to approximately 9. While a pH of 11 would assist with theprecipitation of carbonates the pH of 9 at the higher temperaturewill not assist with the dissolution of Mg from serpentine.

Figs. 11–13 shows our own titration curves for tripropylamine,tributylamine and tripentylamine at various temperatures. For tri-propylamine, the final pH is approximately 9.5 at 18 �C and 7.1 at85 �C. For tributylamine, the final pH is approximately 8.6 at 5 �Cand 6.0 at 85 �C. For tripentylamine, the final pH is approximately6.5 at 5 �C and 4.0 at 85 �C. It follows that with a temperature risefrom 5 to 85 �C, the change in final pH is approximately 2.5 pHunits. This work shows that at elevated temperatures an acidloaded alkylamine with a long chain length behaves similarly toa weak acid and therefore might have the ability to dissolve Mgfrom serpentine.

The equivalence point (point of highest gradient) was found foreach titration curve from which pKa values were estimated. Table 2shows results from this analysis and Fig. 14 shows the constantsfor tributylamine and tripentylamine as a function of temperature.Lines of best fit have been drawn through the data and extended to135 �C. Fig. 14 also shows a line corresponding to the pKa valuesreported by Hamborg and Versteeg [18] for triethylamine. Thelines obtained from this work appear to decrease more steeplythan those for triethylamine, which may be due to experimentalerror associated with vapour losses from the system which couldconcentrate the protons and give lower pH values.

Further experiments to obtain more accurate titration dataincluding data at higher temperatures and pressures (100–150 �Cand 1–5 bar) is planned for the future. The experiments at highertemperature will also involve treating serpentine with the regener-ated amine and acid mix to study the dissolution behaviour. Theseexperiments are akin to treating serpentine with weak organicacids. There is currently a lack of studies on the behaviour of ser-pentine with weak organic acids particularly at high temperatureswhere the kinetics of dissolution is favourable. Teir et al. [19] havereported the behaviour of carboxylic acids alongside stronger acidshowever the studies were confined to 20 �C. Unlike strong acidswhich provide a high initial concentration of protons which de-creases as the mineral dissolves, weak acids provide a low concen-tration that is maintained as the mineral dissolves, and protons areconsumed in the acid-base reaction thereby driving the dissocia-tion of more protons from the weak acid.

K.M. Steel et al. / Fuel 111 (2013) 40–47 47

If the regenerated amines enable high degrees of serpentine dis-solution, experiments will move to examining the extent that theamines can be continually recycled.

4. Conclusion

The best conditions for the dissolution of Mg from serpentinehave been found to be a solids to solution volume of more than50 g/L to enable a high proton concentration. The amount of acidshould be no more than twice the stoichiometric amount neededfor reaction. Reaction temperature should be as high as possible,close to the boiling temperature of the solution or higher (100–150 �C) if using a pressurised vessel. These conditions combinedwith a residence time of 3 h are able to dissolve approximately85% of the Mg in serpentine.

These experiments have shown that the pH needed for magne-sium carbonate precipitation is approximately 8.2 and that trieth-ylamine and tripropylamine are capable of enabling this at 18 �C. Itappears that an excess of amine is needed to maintain the pH of 8while CO2 is bubbled through the solution. So far yields of around20–40 wt.% have been achieved for tripropylamine using residencetimes of approximately 1 h. Precipitation occurred more rapidly fortriethylamine owing to the higher pH generated.

The association of tertiary amines with HCl has been found todecrease with increasing temperature such that there is a differ-ence of approximately 2.5 pH units between 5 and 85 �C. Thismeans that an amine capable of achieving a pH of 8.2 at low tem-perature generates a pH of 5.7 in solution when heated to 85 �C.While this is not low enough to provide a high rate of serpentinedissolution it is thought that increasing the temperature beyond85 �C may yield pH levels capable of dissolving high levels of Mg,particularly given that high temperatures aid the kinetics of disso-lution. Further work is required to study the dissolution behaviourof serpentine with regenerated amine solutions at elevated tem-peratures and pressures.

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